Environ. Sci. Technol. I990,24,363-373
Gagosian, R.; Smith, S. 0.;Nigrelli, G. E. Geochim. Cosmochim. Acta 1982,46,1163-1172. Lefebvre, G.; Germain, P.; Schneider, F. Bull. SOC. Chim. Fr. 1980,1-2, 11-96-11-97, Stohs, S. J.; Haggerty, J. A. Phytochemistry 1973, 12, 2869-2872.
(40) Fauve, A.; Kergomard, A. Tetrahedron 1981,37,899-901.
Received for review March 20,1989. Accepted November 9,1989. Financial support from EROS-2000 Project is gratefully acknowledged. W e thank J.I. Gomez-Belinchon, R. Llop, and M. Valls for technical assistance.
Oxidation of Aniline and Other Primary Aromatic Amines by Manganese Dioxide Shonali Laha and Rlchard G. Luthy" Department of Civil Engineering, Carnegie Mellon University, Pittsburgh, Pennsylvania 152 13 This investigation evaluated the redox reaction between a manganese dioxide, 6-Mn02, and anilines and other aromatic reductants in aqueous suspensions at pH values ranging from 3.7 to 6.5. The reaction with manganese dioxide may represent a pathway for transformation of aniline and other primary aromatic amines in acidic mineralogical and soil/water environments in the absence of oxygen and substantial microbial activity. The reaction rate with aniline is pH-dependent, increasing with decreasing pH, and first order with respect to 6-Mn02and organic solute. Aniline and p-toluidine are demonstrated to be 2-equiv reductants, as is believed to be the case for the other aromatic solutes considered in this study, including the substituted anilines, and hydroquinone and catechol and their alkyl substituents. Ring-bound nitrogen-containing aromatic solutes (methylimidazole,quinoline, and 5,5-dimethylhydantoin) were unreactive with manganese dioxide at pH 6.4. The order of the reactivity of para-substituted anilines was methoxy >> methyl > chloro > carboxy >> nitro; the relative reactivity of these compounds correlated with the solute's half-wave potential and Hammett constant. The principal oxidation products of aniline and p-toluidine with manganese dioxide at pH 4 were azobenzene and 4,4'-dimethylazobenzene, respectively, which agreed with a postulated oxidative-coupling reaction mechanism. The abiotic redox reactions of primary aromatic amines and azo compounds may result in various respective oxidative-coupling and reductive-decoupling reactions. These processes may be significant with regard to the persistance and transformation of these classes of organic contaminants in environmental systems.
Introduction The purpose of this investigation was to assess the rate of the redox reaction between a manganese dioxide, 6Mn02(s),and various aromatic reductants including several nitrogen-containing compounds. The initial rate and order of the reaction with respect to the reductant was determined as well as the effects of manganese dioxide concentration and pH. The principal reaction products, under acidic conditions, resulting from the oxidation of aniline and p-toluidine were determined. In natural waters, Mn(II1) and Mn(1V) are usually present in the form of sparingly soluble oxides and hydroxides, whereas Mn(I1) is the soluble phase. In soil/ sediment environments, manganese oxide is believed to be among the strongest oxidizing agents that may be encountered in the absence of molecular oxygen. Manganese oxide can be reduced and dissolved by organic compounds, increasing the mobility of manganese and its availability to organisms ( 1 , 2 ) . These oxidative processes involving manganese oxides may constitute an important abiotic degradative pathway for organic compounds in subsurface 00 13-936X/90/0924-0363$02.50/0
environments. Earlier studies by Stone and Morgan (2) demonstrated some features of the reductive dissolution reaction between manganese oxide and organic solutes and some of the factors that influence the rate of the reaction. Stone (3) considered the reductive dissolution of manganese(III/IV) oxides by substituted phenols. Those investigations were performed with various manganese oxide suspensions, one of which was primarily the mineral-phase feitknechtite, P-MnOOH(s), with some amount of manganite, y-MnOOH(s). In this study, a hydrous manganese dioxide suspension was prepared according to Murray (4), for which his stoichiometric and X-ray diffraction analyses indicated that this synthetic manganese dioxide is structurally similar to the naturally occurring mineral birnessite, 6-Mn02(s). The organic reductants investigated in this study were aniline and various substituted anilines, hydroquinone and catechol and some of their alkyl substituents, and several ring-substituted nitrogen-containing aromatics. Figure 1 shows the structures of the compounds discussed in this study. Anilines and other aromatic amines may originate as environmental contaminants from the use of pesticides and herbicides, as well as from chemical manufacturing residues, and from byproducts of energy technologies. Along with phenol, aniline is listed as a high-priority compound in the study of pollutants from coal-conversion process wastes (5). Aniline residues are formed in the soil as a result of microbial and plant metabolism of phenylurea, acylanilide, phenylcarbamate, and nitroaniline herbicides (6). Chlorinated anilines such as 2,4,5-trichloroaniline, 4-chloroaniline, 3,4-dichloroaniline, and 2,6-diethylaniline may be released as degradation products and intermediates of various phenylurea and phenylcarbamate pesticides (7, 36). Aniline derivatives occur as intermediates in dye-stuff manufacture and this constitutes another possible source of environmental aniline contamination. A number of substituted anilines may be carcinogenic (8). Aniline and toluidine, i.e., methylaniline, and other aromatic amines are generally toxic and can induce various adverse physiological responses (8, 9). For these reasons it is of considerable interest to explore the physical, chemical, and microbial transformations that may alter the toxicity, mobility, and bioavailability of aniline and related compounds in the environment. Aniline and other aromatic amines are subject to complex environmental transformations. Lyons, Katz, and Bartha (6, 10, 11) performed studies on the microbial pathways for aniline elimination from aquatic environments, from which they concluded that biodegradation may be the most significant mechanism for the removal of aniline from pond water. Hwang and Lee (7) concluded that photochemical processes were primarily responsible for mineralization of 2,4,5-trichloroaniline in surface water
@ 1990 American Chemical Society
Environ. Sci. Technol., Vol. 24, No. 3, 1990
363
I
OCH, p-Anisidine
Aniline
CH3 p-Toluidine
c1 p-Chloroaniline
NHz
OH
I
COOH p-Aminobenzoic Acid OH
NO2
p-Nltroaniline
Catechol
CH3 4-Methylcatechol
OH
OH Hydroquinone
OH Methyl hydroquinone
OH
OH 2,3-Dimethylhydroquinone
OH
OH
OH Trimethylhydroquinone
OH tert-Butylhydroquinone
Quinoline
CH3 Methylimidazole
5,5-Dimethylhydantoin
Flgure 1. Structure of the anilines and other aromatic compounds employed in this study.
from an eutrophic lake. Zepp and Schlotzhauer (12) showed that the photolytic degradation of aniline is enhanced in the presence of algae. The principal degradative processes for aniline and other primary aromatic amines in subsurface environments are either microbially induced reactions or abiotic degradation processes through redox reaction with soil constituents, for which the oxidative degradation of aniline by manganese dioxide is reported here. Materials and Methods The inorganic reagents were analytical grade (Fisher Chemicals) and the organic compounds were used as received (purity >99%, Aldrich). All stock solutions of organic substrates were used within 24 h to minimize the possibility of degradation by photolysis or oxidation by oxygen. Deionized water was used in the experiments and all glassware was soaked in 1:5 HN03/water and thoroughly rinsed with deionized water prior to use. Other reagent solutions were also made up in acid-cleaned glassware with deionized water; these solutions were discarded within 2 or 3 days to prevent microbial activity. Reagents and reaction solutions were foil-wrapped,and the more reactive nitrogen aromatics were maintained under inert atmosphere. Preparation of Oxide Suspension. The preparation of approximately 0.01 M b-MnOz(s) was performed as described by Murray ( 4 ) . Murray's analysis of the manganese dioxide produced by his method indicated it to be structurally similar to the naturally occurring mineral birnessite, with a surface area of 263 f 5 m2/g by the BET method and a maximum pH (zpc) of 2.4 by extrapolating electrophoretic mobilities. The manganese dioxide used in this study may be similar to that described by Murray, as the same method of preparation was employed. The procedures for preparation of the manganese dioxide were 364
Environ. Sci. Technol., Vol. 24, No. 3, 1990
as follows: 1. Combine 40 mL of 0.1 M KMn04 and 80 mL of 0.1 M NaOH with stirring; adjust volume to 500 mL. 2. Add 60 mL of 0.1 M MnC1,. 3. Allow 6-MnOZsuspension to settle out. 4. Decant supernatant; wash oxide several times and adjust to 1-L volume with water having the same ionic strength and buffer as to be utilized in the experiments. The reactivity of a manganese oxide suspension may change with time (1,2). To account for the effect of aging, all suspensions were used 24 h after preparation. Tests with Catechols and Hydroquinones. Hydroquinone, methylhydroquinone, 2,3-dimethylhydroquinone, trimethylhydroquinone, tert-butylhydroquinone, catechol, and 4-methylcatechol were evaluated in dissolution reactions with manganese dioxide. The reaction rate was monitored by increase in Mn(I1) with atomic adsorption spectrometry (Perkin-Elmer Model 703). The reaction was performed with continuous stirring under nitrogen purge with no oxygen. Samples of 5 mL were withdrawn by syringe a t predetermined time intervals and expressed through a syringe-tight membrane filter (0.22 wm pore, 1-in. diameter; Millipore Millex-GS). It was verified that the filtration procedure stopped the reaction. The aliquots were acidified with 4 N HN03 and analyzed for dissolved manganese, [Mn(II)ldisS.Where the reaction proceeded very fast, the reacting solutions were flash mixed in a beaker, and samples were withdrawn by syringe immediately and then filtered. The reaction time included the residence time in the syringe, and filtration marked the end of the time interval. The minimum reaction interval by this procedure was about 12-15 s, which represents a practical limit of such a technique. The [Mn(II)]d, is an operational definition of dissolved manganese as employed earlier in Stone and Morgan's work ( I ) . It is that amount of manganese not retained by the membrane filter; and an increase in manganese beyond any that may be present in a reaction blank is taken to mean soluble manganese, Mn(I1). The pH of the reaction mixture was maintained at 6.4-6.5 for these tests with a 0.01 M sodium bicarbonate buffer. Ionic strength was kept at 0.1 M NaN03. Blanks were prepared without organic solute to assess any changes in oxide and Mn(I1) concentrations due to reasons other than reductive dissolution by reaction with the organic solute. The Mn(I1) reported represents net manganese oxide dissolved, accounting for any manganese that may pass the filter initially. Thus, a value of zero is taken for the initial concentration of Mn(I1) produced by reaction. Generally blanks were found not to have significant dissolved manganese. Initial reaction rates were determined from the slope of the manganese dissolution curve by using the first few data points for which the reactant concentrations did not change significantly. Tests with Aniline and Substituted Anilines. The test protocols were similar to those for the hydroquinones, but in addition to evaluating the reaction rate by increase in Mn(II), the organic concentration was also monitored by use of high-pressure liquid chromtographic (HPLC) techniques. Tests were performed at pH 4 with an acetate buffer as well as at pH 6 with a bicarbonate buffer. Chromatographic analyses were performed with a Perkin-Elmer Series 3 liquid chromatograph equipped with a Rheodyne 7105 injector and a Perkin-Elmer LC-15 UV detector at wavelength 254 nm. The column used was a 15-cm Supelco LC-PAH column with a particle size of 5 pm. The solvent used was 50:50 acetonitrile/water, at a flow rate of 1 mL/min. Injection volumes were between 4 and 9 pL. Where higher concentrations of reactants were
used the buffer concentration was also increased, e.g., 0.1 M acetate for 0.005 M reactants at pH 4, and 0.1 M bicarbonate for pH values of 6. Samples were withdrawn at predetermined time intervals by syringe and then filtered (0.2-pm pore, Millipore Millex-FG Teflon filters) into 5-mL vials that were sealed, covered with foil, and refrigerated until analyzed by HPLC techniques. Representative aqueous stock solutions of organics were analyzed both before and after filtration to ensure lack of organic adsorption on the Teflon filters. Reaction blanks were also studied to observe any change in aniline concentration in the absence of both oxygen and 6-Mn02. Aniline degradation was negligible under such conditions in the time period of the experiment. Peak areas for the organic analysis from HPLC determinations were divided by the quantity injected to yield response factors from which calibration curves were constructed with known solute concentrations. Graphs of solute concentration remaining versus time were plotted and the initial slopes were used to determine the initial reaction rates. Identification of Reaction Products. The principal reaction products for aniline and p-toluidine were determined by GC-MS techniques. These samples were prepared for reaction by mixing 0.01 M organic and 0.01 M 6-Mn02(s),a t pH 4 with 0.1 M acetate buffer and 0.1 M NaN03ionic strength, in a separatory flask. The flask was shaken intermittently in a gentle manner to ensure proper mixing. L-Ascorbic acid (1 M) was added to stop the reaction by dissolving the remaining manganese oxide (1). The dissolution of residual manganese oxide also provided for release of any adsorbed reaction products. The pH of the mixture was then adjusted to 12 by the addition of sodium hydroxide pellets, and the organic compounds were extracted with methylene chloride by shaking vigorously for 10 min, allowing the CH2C12to separate, and withdrawing through a Teflon stopcock. The methylene chloride extract was concentrated by use of a KudernaDanish evaporator concentrator to volatilize the solvent. In these tests filtration was avoided in the event that a reaction product may have been retained on the filters. The concentrated extract was injected into a 5985 Hewlet-Packard gas chromatograph-mass spectrophotometer with a 25 m X 0.2 mm SB-Smectic column. The injection temperature was 280 "C, the carrier gas helium at 1 mL/min, the ionization potential 70 eV, and the voltage 2400 V. The program was 1min at 40 "C, 40-220 "C at 20 "C/min and 220-280 "C at 4 OC/min. The mass spectra indexes used were obtained from the literature (13, 14). Standards were also run under the same conditions to confirm the identity of the reactants and major products. The confirmation was done by matching the chromatographic retention time and comparing the principal mass ion distribution of standards and samples with those reported in the compendia of mass spectra. Stock solutions of aniline and p-toluidine were analyzed by GC-MS to ensure purity and the absence of possible degradation products. Control samples having organic solute and buffer, but no 6-Mn02(s),were carried through the mixing and extraction procedure, including the addition of ascorbic acid to verify the lack of reaction products and to demonstrate that identified organic products were present as a result of redox reaction with 6-Mn02(s), rather than as a spurious experimental artifact. Additional Tests. Rate tests were performed with the following ring-substituted nitrogen-containing organic compounds: methylimidazole, quinoline, and 5,5-dimethylhydantoin. The pH was maintained a t 6-6.4 and
--- Experimental Data -Estimate of Initial Reaction Rote
.-
,,*x , 6
1 ,/1/'
Initial Rate= 4.5x10-5mo1es/~-min
' I
tl
1I/
G C
z -
pH =4.4
[S-MnO ],=6.8x IO+ M [AnilineIo = 2 x I O + M
0 0
10
20
Time, minutes
Figure 2. Reductive dissolution of manganese dioxide by aniline at pH 4.4.
the reaction was monitored by the concentration of dissolved manganese (using atomic absorption spectrometry techniques) over a period of days.
Results Initial Reaction Rates with Aniline. Rate tests were performed to evaluate the dissolution of manganese dioxide by several nitrogen-containing aromatic compounds. The compounds considered were aniline, p-anisidine, quinoline, methylimidazole, and 5,5-dimethylhydantoin. The objective was to observe the rate of dissolution of manganese dioxide by representative nitrogen-containing aromatics in order to identify those compounds that are reactive with manganese dioxide. The reaction was monitored by increase in the concentration of dissolved manganese. At pH 6.4 and millimolar concentrations of organic and oxide, dissolution of 6-Mn02(s)by methylimidazole was extremely slow, and unambiguous results were not obtained over 1 month. Similarly, quinoline and 5,5-dimethylhydantoin did not exhibit significant dissolution rates under these test conditions. The initial rate with aniline and 6-Mn02(s)was relatively slow at a pH value of 6.5, and the reaction rate increased with decreasing pH. Figure 2 shows that at pH 4.4 dissolution of manganese dioxide by aniline gave an initial reaction rate of 4.5 X mol/L.min expressed in terms of rate of dissolution of manganese, and that the reaction appeared to proceed a t a slower rate after -20 min. Subsequent experiments with aniline were monitored by determining the concentration of aniline remaining in solution with HPLC techniques, for which a typical result gave an initial reaction rate of 8 X mol/L.min expressed in terms of consumption of aniline for test conditions similar to those shown in Figure 2, except for a small proportional increase in the initial aniline concentration. The similarity of the initial reaction rates under comparable test conditions as given by either manganese dissolution or consumption of aniline, in conjunction with identification of aniline reaction products as discussed later, suggests an equivalent stoichiometry in the reduction of aniline by manganese dioxide. An analogous conclusion regarding stoichiometry was obtained by Stone and Morgan (1,2) for the case of reductive dissolution of pMnOOH(s) by hydroquinone, which was applied in their work for the case of 15 aromatic and 12 aliphatic comEnviron. Sci. Technol., Vol. 24, No. 3, 1990
365
-3
/-
/
' 1 -3.5
.-
/
E
9
slopen=I
pH.4
-
:/
.c
.+ Y Cn
-
{
log [Ar~iline]~ , moles/J}
Flgure 3. Estimation of the order of initial reaction rate with respect to aniline concentration for oxidation of aniline by manganese dioxide.
pounds behaving as 2-equiv reductants with manganese oxide. A similar inference regarding the compounds reacting as 2-equiv reductants was used in this study to assess the rate of the reaction for 6-Mn02(s)with aniline and substituted anilines, as well as for catechols and hydroquinones. Hence, the amount of organic compounds remaining could be expressed in terms of dissolution of manganese: [organic], = [organiclo- [Mn(II)],
Environ. Sci. Technol., Vol. 24, No. 3, 1990
-3
-2
-2.5
log{ [8-Mn02]o, moles/Q}
Figure 4. Estimation of the order of initial reaction rate with respect to manganese dioxide concentration for reduction of 6-Mn0, by aniline.
(1)
where [organic], represents the molar concentration of organic solute a t time t, and [Mn(II)], represents the difference in molar concentration of dissolved manganese between the reaction mixture and blanks at time t. While subsequent identification of principal reaction products in acidic samples suggests that aniline and p-toluidine are primarily 2-equiv reductants, additional work is needed to determine if the anilines are exactly 2-equiv reductants under various conditions. Reaction Order with Aniline. The reaction kinetics of aniline and manganese dioxide were investigated to determine the order of the reaction with respect to the initial organic solute and oxide concentrations and the dependence of the initial reaction rate on the pH of the reaction mixture. The order of the initial reaction rate with respect to aniline was evaluated by tests in which the oxide concentration was employed at a relatively high value of 5 X M, while maintaining constant pH and buffer strengths through comparative tests and varying the aniline concentration by an order of magnitude from 0.5 X to 5 X M. In this approach, the initial reaction rate is computed over the first few time increments of the reaction, for which it may be assumed that the change in concentration of the reactants is negligible. The results shown in Figure 3 demonstrate that the reaction rate is first order with respect to aniline. Similarly, the order of the reaction with respect to 6Mn02(s) was determined by varying the initial oxide concentration from 1.25 X to 5 X M, while maintaining the initial aniline concentration at 2.5 X lom3 M and pH at 4. Figure 4 indicates that the reaction order with respect to manganese dioxide concentration is approximately unity. The effect of pH on the initial reaction rate was determined a t six different pH values from 4.0 to 6.5 at an initial aniline concentration of 2 X M and an initial manganese oxide concentration, [6-Mn02]o,of 5 X M. The results are shown in Figure 5, from which it is seen 366
-5
c
Cn
0
-6
'
4
I
I
5
I
I
I
6
PH
Figure 5. Effect of pH on the initial reaction rate for oxidation/reduction reaction between aniline and manganese dioxide.
that pH has a marked effect on reaction rate with the rate decreasing by an order of magnitude over this pH range. The reaction order is not constant with respect to pH, increasing from -0.3 at pH 4 to -0.6 at pH 6.5. Although pH values generally encountered in natural systems are higher, a pH of 4 was chosen in subsequent tests to assess reaction products and to evaluate the effect of substituents on the reaction rate of anilines. This was done for purposes of experimental convenience because of the ease of generation of data in a short period of time; also the blank values were negligible the shorter the reaction time considered. It is appreciated, however, that the reaction rate may diminish substantially at higher pH, as suggested by data in Figure 5. Substituent Effect and Reaction Order for Hydroquinone and Catechol. Hydroquinines and catechols are benzyl compounds with two hydroxy substituents, at para and ortho positions, respectively. A study of the dissolution of manganese was performed with hydroquinone, catechol, and several alkyl substituents: methylhydroquinone, 2,3-dimethylhydroquinone,trimethylhydroquinone, tert-butylhydroquinone, and 4-methylcatechol. The initial concentrations of the organic solutes to 5.0 X M. were varied from 0.5 X 6-Mn02 was used in these tests at (1.7-2.3) X M concentration with pH maintained at 6.5 by using 0.01 M
Table I. Initial Reaction Rates for Hydroquinone and Alkyl-Substituted Hydroquinones at pH 6.5 init concn, mol/L [6-MnO*Io [ organiclo 1.7 x 1.7 x 2.3 x 1.9 x a
0.5 x 1.5 x 2.5 x 5.0 x
10-3 10-3 10-3 10-3
init reaction rate: mol/L.min methyl-HQ 2,3-dimethyl-HQ
HQ 2.4 X 1.5 x 1.5 x 2.3 x
10-3 10-3 10-3 10-3
lo4 10-3 10-3 10-3
2.1 1.4 2.0 2.2
x 10-4 x 10-3
tert-butyl-HQ
2.0 x 10-4 5.2 x 10-4 1.0x 10-3 1.9 x 10-3
x 10-3 x 10-3
1.6 X 5.2 x 1.0x 1.9 x
trimethyl-HQ
lo4 10-4 10-3 10-3
1.4 X 4.4 x 5.9 x 1.8 x
10” 10-4 10-4 10-3
HQ, hydroquinone. -6.3
pH.6.5 [8-MnO2],= 1 . 7 IOm3 ~ M
-6.4
2 , 3 -dimethyl hydroquinone
[ 0rgonic]~=2.5x I O - ~ M
h -I
>
-W
-6.5
-
7J//
trimethyl hydroquinone
E
-d
-6.6
F
u ca *
-6.7
C
9
-6.8 -4
-6.9
I
-3
{
log [~rganic]~ ,moies/l}
2
Time,minutes
Figure 6. Reductive dissolution of manganese dioxide by hydroquinone and alkylsubstituted hydroquinones at pH 6.5.
NaHCO, buffer. Dissolution of 6-Mn02was measured over time. The results were expressed as the natural logarithm of the oxide remaining in the reaction mixture versus time since the start of the reaction, for which a typical set of results obtained for the hydroquinones is shown in Figure 6. These data indicate that the reaction may slow with time, and thus only the initial reaction rates were considered for comparison, as summarized in Table I. It is observed that reaction rate increases with the initial concentration of organic reductant, and results shown in Figure 7 for dimethylhydroquinone and trimethylhydroquinone indicate that the reaction order with respect to initial organic loading is close to unity. This is consistent with the data presented in Table I, for which the initial reaction rates for the hydroquinones increase by approximately an order of magnitude as initial organic concentration increases from 5 X to 5 X M. Methylhydroquinone and hydroquinone are the most reactive among the five compounds. However, the substituent effect on reaction rate for the hydroquinones is not particularly significant. As the reaction proceeds, the data in Figure 6 show deviation from linearity, suggesting that the reaction processes are more complex than implied by comparison of first-order initial reaction rates. Results in Table I1 show the initial reaction rates for catechol and 4-methylcatechol. The initial reaction rates are in the range of those observed for the hydroquinones. As with the hydroquinones, it was observed that the dissolution reaction with catechol was rapid, with the reaction slowing appreciably after 2 min for these test conditions. The initial reaction rates increase for 4-methylcatechol in proportion to the organic concentration. However, a t a M organic the reaction rate for concentration of 5 x both catechol and 4-methylcatechol is no different from
-2
Figure 7. Estimation of the order of the redox reaction with respect to 2,3-dimethylhydroquinone and trimethylhydroqulnone for reduction of manganese dioxide.
Table 11. Initial Reaction Rates for Catechol and 4-Methylcatechol at pH 6.4 init concn, mol/L [6-MnO2Io [organic],, 1.7 x 1.9 x 1.7 x 2.3 x
10-3 10-3 10-3 10-3
0.5 x 1.5 x 2.5 x 5.0 x
10-3 10-3 10-3 10-3
init reaction rate, mol/L.min catechol 4-methylcatechol 1.0x 10-3
1.5 x 10-3 2.6 x 10-3 2.7 x 10-3
4.5 x 1.7 x 2.7 x 2.7 x
10-4 10-3 10-3 10-3
that at organic concentration 2.5 X lo-, M. Stone (3) reported similar behavior for the case of p-methylphenol and manganese(III/IV) oxide in which the reaction rate begins to level out as the pH is decreased below 5.0. This is explained through a reaction scheme in which precursor complex formation and electron transfer are rate-limiting steps and electron back-transfer is negligible. In that case, the reaction rate resembles a transition from pseudo first order to zero order with respect to the concentration of the organic compound. In Stone’s work it was explained that, at lower concentrations of p-methylphenol, the reaction rate increases linearly with organic concentration, but as the concentration of the organic compound increases, a point is reached where the oxide surface is saturated with organic compound and electron transfer becomes rate limiting, and the reaction rate remains unaffected by higher concentrations of organic. In this study the results for catechol and 4-methylcatechol indicate that as the initial concentration of the organic solute increases beyond the initial concentration of 6-Mn02, the proportionality between reaction rate and organic concentration no longer holds, probably through a change in rate-limiting mechanism as suggested by Stone (3). The formation of Mn(I1)-organic complexes was not investigated. Stability constants reported for catechol by Environ. Sci. Technol., Vol. 24, No. 3, 1990 367
Table 111. Initial Reaction Rates for Aniline and Substituted Anilines in the Presence of 6-Mn02 compound p-anisidine p-toluidine aniline p-chloroaniline aminobenzoic acid p-nitroaniline
EljZ?
v
Hammettb const
[organiclo, mol/L
[MnOzlo, mol/L
p~
init rate, mol/L.min
kex L/moiimin
-0.78 -0.31
6.2 X 2 x 10-3 2.5 x 10-3 5 x 10-3 5 x 10-3 2 x 10-3
3.3 x 10-5 2.5 x 10-3 5 x 10-3 5 x 10-3 5 x 10-3 5 x 10-3
4.4 4 4 4 4 4
6.5 X lo4 1.2 x 10-3 1.3 x 10-4 1.1 x 10-3 3.1 x 10-3 -1 x 10-7
32000 240 10.4 44 3.44 -0.01
0.393 0.537 0.625 0.675 0.794 0.935
0.11 0.42 0.79
?
"E,,, values from Suatoni et al. (28). u+ values from March (22).
54x'65
t //'
/'e
I n i t i a l Rate = 6 5 x 10-smoles/l-min
1
pH=4,Time?O mins
pH.4.4
- M no,] 0=3. 25 x I 0-4M
5
I0
I5
20
25
30
Reaction Blank
[ p - A n i ~ i d m e ] ~ =26 5 ~ 1 M0 ~ ~
", 5
I0
I
I5
20
25
30
Time, minutes
---
Experimental D a t a Initial Reaction Rate
-Estimate of
Figure 9. Chromatogram from GC-MS analysis for identification of reaction product for 10-min redox reaction of aniline and manganese dioxide.
I
0
50
IO0
T i m e , seconds
stants, kexp, are also presented in the table. The rate constants were computed as
Figure 8. Reaction of p-anisidine and manganese dioxide at pH 4.4.
kexp = (initial reaction rate)/[6-MnOz]o[organic]o
Sillen and Martell (37), e.g., log K1 = 7.52, suggest that Mn(I1)-organic complex formation in the aqueous phase may have an effect on the reaction mechanism and rate. Smith and Martell (38)reported no stability constants for Mn(I1) with aniline, p-toluidine, or p-anisidine. However, reported log K values are small for Cd(I1) with aniline (log K = O.l), p-toluidine (log K = 0.261, and p-anisidine (log K = 0.45), and by inference, complexation of the anilines with Mn(I1) may be insignificant in comparison to catechol. Substituent Effect for Aniline. The initial reaction rate was evaluated for various substituents of aniline: p-anisidine (i.e,, p-methoxyaniline), p-toluidine &e., pmethylaniline), p-nitroaniline, 4-aminobenzoic acid, and p-chloroaniline. These tests were performed at pH 4 with acetate buffer. An initial oxide concentration of 5 X M was employed except for the tests with p-toluidine and p-anisidine, for which [6-Mn02]o= 2.5 X and 3.25 X lo4 M, respectively. The organic solute concentration was (2-5) X M, except for tests with p-anisidine, for which it was 6.25 X M. p-Anisidine dissolved manganese oxide so rapidly that it was not possible to measure the dissolution rate at millimolar concentrations of both oxide and organic. Figure 8 shows a test result for increase in Mn(I1) for p-anisidine where the initial concentrations of and 3.25 X M, organic and 6-Mn02 are 6.25 X respectively. In this case the reaction was complete in -2 min, as the amount of manganese dissolved approached the stoichiometric amount of p-anisidine. For the tests with the other compounds, the rate was substantially slower and the course of the reaction was assessed by disappearance of reductant. Table I11 summarizes the initial reaction rates of the various substituted anilines. The experimental rate con-
where psuedo-first-order dependence has been assumed with respect to both organic compound and manganese dioxide for the initial reaction rate as discussed earlier. p-Nitroaniline was found to be particularly unreactive; after a reaction period of 1 2 h there was only -5% disappearance of the compound at pH 4. As with p-anisidine, p-toluidine reacted readily with visible color production; the oxide suspension became wine-red on addition of organic, and dissolution was pronounced. p-Chloroaniline and 4-aminobenzoic acid also reacted rapidly with visible color changes, the product being a deep peach for the former and orange for the latter. Reaction Products. The identity of the reaction products was investigated for the reductive dissolution of manganese dioxide by aniline and p-toluidine. A set of tests with aniline was conducted by mixing 5 mL of M aniline with 5 mL of M 6-Mn02in acetate buffer at pH 4 in a centrifuge tube. The reaction was stopped at 10 min by the addition of 30 mL of 0.05 M L-ascorbic acid, which dissolved the remaining oxide suspension immediately. Aside from stopping the reaction, the purpose of this step was to release any reaction product that may have adsorbed on unreacted manganese dioxide and to prevent problems of solids separation and emulsification during subsequent extraction of the reaction mixture. The extraction was performed by adjusting the pH of the reaction solution t~ 12 with sodium hydroxide and extracting with 10 mL of methylene chloride for 10 min with vigorous shaking. The extract was concentrated by a factor of 13 and the concentrated extract was analyzed by GC-MS procedures. A typical chromatographic result is presented in Figure 9, which shows the presence of aniline at retention time 4.3 min and azobenzene at retention time 19.8
368
Environ Sci Technol, Vol 24, No 3, 1990
(2)
was still the principal reaction product, although a number of minor species were apparent also. p-Aminodiphenylamine was identified as one of the minor reaction products. Azobenzene was confirmed by comparison of retention time and spectral intensities of an analytical standard. Figure 11 shows an inset of mass spectra for the reaction products corresponding to Figures 9 and 10 compared with a mass spectrum of an azobenzene standard. The four largest intensities for the azobenzene standard are m/z 77, 182,105, and 152, which agree with the spectral displays for the identification of the compound as reaction product and which are listed as intense peaks in mass spectra indexes (13, 14). Figure 1 2 shows a chromatogram of reaction products for 5 X equimolar concentrations of p-toluidine and 6-Mn02 for a 10-min reaction time at pH 4. The predominant reaction product is clearly 4,4'-dimethylazobenzene, although several unidentified species were also present as minor reaction products. GC-MS determination of 4,4'-dimethylazobenzene was confirmed with an analytical standard by comparison of the chromatographic retention times and the largest spectral intensities, i.e., mass ion distribution of m/z 91, 210, and 119.
Reaction of Aniline with 8-Mn02 pHz4,Time 30 mins
5
10
20
15
25
Time, minutes
Figure 10. Chromatogram from GC-MS analysis for identification of reaction products for 30-min reaction of aniline and manganese dioxide showing azobenzene as the principal product with various minor products including 4-aminodiphenyiamine.
min. The chromatogram shows no apparent reaction products other than azobenzene. Sample blanks were prepared for analysis in the same manner with the omission of the manganese dioxide, in which the manganese suspension volume was replaced by distilled water. Figure 9 also shows the chromatogram for the reaction blank, which indicates that sample manipulation and preparation for analysis, including the addition of ascorbic acid, resulted in no reaction product formation, confirming that the azobenzene is a product of the reaction of aniline with manganese dioxide, rather than being attributable to other degradative processes or impurities in the aniline. Figure 10 shows a chromatogram of products for reaction of 6-Mn02and aniline at equimolar concentrations of 5 X in a separatory flask in which the reaction was allowed to proceed for 30 min prior to analysis. In this test, larger volumes (100 mL) of reactants were used so that the subsequent extract would result in a greater concentration factor. In this case of the longer reaction time, azobenzene 100 -. 80 - .- 6 0 4020+ c 0H
.z5 ~
$ 2
Azobenzene Standard
,
loo-. 806040-
20 -. 0
Discussion For the reaction times and conditions considered in this investigation, degradation of aniline is negligible in the absence of manganese dioxide with exclusion of oxygen. This is consistent with the observations of Lyons et al. (11) who reported that autoxidation of aniline in the absence of light contributes little to aniline removal. The dissolution rate of manganese dioxide by aromatic amine compounds depends on the nature of the oxidant, the reactant concentrations, and the medium composition. pH is seen to play an important role since the dissolution rate depends on the redox potential and on the extent and nature of the interaction of the aromatic solute with the oxide surface. The apparent reaction order with respect
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Azobenzene Reaction Time: IOmin
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Azobenzene
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Environ. Sci. Technoi., Voi. 24, No. 3, 1990
389
5 4,4'- Dimethytarobenzene
p- Anisidine
/
\
\
4
\
\\ \
Reaction of p-Toluidine with 8-MnOp pH.4,Time IO mtns
3
\
\
\
0' \\p-Chlorophenol
p- To I u i di n e P-MethYl Phenol
2
@ Phenolb$-Chloroaniline
a W
5
10
r m
15
Time,minutes
Flgure 12. Chromatogram from analysis of reaction products for 10-min reaction of p-toluidine and manganese dioxide. 4,4'-Dimethylazobenzene Is the prlncipal reaction product.
to [H+]is derived from the slope of the logarithm of the initial reaction rates plotted as a function of pH: initial rate = k[organic]'[manganese oxide]"[H+]"
AnilineO \,
I
0 -
p-Hydroxybenzoic acid
The order of this reaction is demonstrated to be unity with respect to the concentration of both manganese dioxide and aniline, and at constant pH the rate law has the form reaction rate = k,[aniline] [manganese dioxide] (6) Rate laws of a similar form have been proposed for reaction of aniline and substituted anilines as 2-equiv reductants in aqueous acetic acid medium for oxidation by peroxydisulfate, Sz082- (15), thallium triacetate, Tl(OAc), (16), and iodate, IO3-(17). The rate constants for the oxidation of the aromatic amines may be correlated with the organic solute half-wave potential and the substituent Hammett constant. Halfwave potentials have been reported (18) from anodic voltammetry experiments with para- and meta-substituted 370
Environ. Sci. Technol., Vol. 24, No. 3, 1990
ocld
\\\ o-Hydroxybenzoate O
0
p-Hydroxyacetophenone @\ \ \ 0 Anilines \ pH=4,8-Mn02
-I
O Phenols
-2
pH.4.4, Mn (III/IV) Oxide
+
The order with respect to [H+],n,need not necessarily be constant or an integer number. Figure 5 showed that reaction rate increases as the pH is decreased for the reaction of 2 X M aniline with 6-Mn02. The apparent order of the reaction with respect to [H+]defines a curve, rather than a straight line, varying in slope from 0.32 to 0.64. Stone (3) observed that the reductive dissolution of Mn(III/IV) oxide by p-methylphenol varied with pH, with n decreasing from about 1.2-1.3 at pH >6 to about 0.44.5 at pH 4. The test results with p-methylphenol, as with aniline, suggested that the observed pH dependence may arise from specific interactions between the reductant and the oxide surface sites, including protonation reactions that promote the formation of surface precursor complexes and/or increases in the protonation level of surface precursor complexes that increase rates of electron transfer (3). The experimental results suggest an equivalent stoichiometry in the initial reaction of aniline with manganese dioxide by evidence of the principal reaction products, Le., Mn(I1) and azobenzene, and by the similarity of the rate of disappearance of the reactants. Aniline is thus a 2-equiv reductant, for which the similarity in initial reaction rate with regard to appearance of Mn(I1) or depletion of aniline demonstrates that d[Mn(II)]/dt = -d[aniline]/dt = reaction rate (5)
0 p-Aminobenzoic \
(3)
or log (initial rate) = log (k) + 1 log [organic] m log [manganese oxide] + n log [H+] (4)
b
\
\
P\
p-Nitroaniline
-3 0
0.2
0.4
0.6
0.8
1.0
E b2, volts Flgure 13. Correlation of half-wave potentials with experimental rate constants for redox reaction of substituted anilines wtth 6-Mn0,, and substituted phenols with manganese(III/IV) oxide. Rate data for the phenols are from Stone (3).
phenols and anilines for use in polarographic identification of oxidizable and reduced organic compounds. The half-wave reduction potentials depend upon the chemical speciation and diffusion coefficients of oxidant and reductant as well as the thermodynamic driving force for the reaction (3, 19). The tendency to donate electrons increases as the solute half-wave potential decreases (18). Logarithms of the experimental rate constants, kexp,of substituted anilines at pH 4 are plotted against half-wave potentials, El 2, in Figure 13. The keXpvalues were presented in $able I11 with reported values of El12.The reaction rate is observed to generally decrease with increasing half-wave potential. p-Anisidine has the lowest half-wave potential and reacts most quickly with manganese dioxide, while p-nitroaniline with the highest halfwave potential is the least reactive. Data from Stone (3) for reaction of the substituted phenols with manganese(III/IV) oxide at pH 4.4in acetate buffer are also presented in Figure 13, and these data agree with the trend observed for substituted anilines using Ellz values for the substituted phenols as reported also by Suatoni et al. (18). The particular manganese oxide employed by Stone (3),for which the rate data are compared in Figure 13, was prepared by reaction of MnSO, and NaMnO, at pH 6.6; this was similar to the manganese dioxide preparation procedure employed in this investigation except that Mn2+and Mn0,- were reacted under alkaline conditions. The chemical conditions are similar for the two sets of tests shown in Figure 13, with pH 4 used in the tests with the substituted anilines and pH 4.4 in the tests with the substituted phenols. These results suggest that the reactivities of 6-Mn02and manganese (III/IV) oxide in acidic environment are of the same magnitude with respect to reductive dissolution of manganese dioxide by either substituted phenols or anilines. Hammett constants are linear free energy parameters that may be used to estimate chemical properties for or-
H-N-H
H-N-H t++,-
Aniline
H-N-H
H-N-H
y-
/
H
Resonance Structures of Cotion Radical which undergo Coupling Reoctions Cation Radical Coupling Product head-to-tail H , N o N H o
+ ZH'
-
Final Product H
N N
0
~ +zH++ 2e-
4 aminodiphenylomine
\
I
-0.8
-0.4
0
-2
m;
+
,
0.8
0.4
benzidine
hydrozobenzene p-Nitroaniline'
Flgure 14. Correlation of Hammett u+ constants with experimental rate constants for redox reaction of substituted anilines wtth 6-Mn0,. Nitroaniline does not follow the correlation, see text for discussion on the behavior of this compound.
ganic solutes (35),such as acid dissociation constant or rate of hydrolysis (201, as well as rates for many other types of reactions of aromatic solutes (21). The estimation method is basically a substituent effect approach in which a correlation is established between the chemical property of the parent compound (e.g., phenol or aniline) and that of the substituted compound. I t has been reported that the Hammett substituent constants, u, are linearly proportional to the change in electron density at the aromatic ring carbon meta or para to the primary substituent group (18). Since the value of the half-wave potential for substituted phenols and anilines depends on the electron density at the carbon atom to which either the hydroxy or amine group is attached, the half-wave potentials of meta- and para-substituted phenols and anilines should be linearly related to Hammett's u constants. The selection of the appropriate set of Hammett constants (u, u,+ or a-) depends on the reaction mechanism within a given series. Hammett u+ values are used where the substituent group can enter into direct resonance with the reaction site in the transition state for an electron-donating group (e.g., NH,) developing a positive charge (21, 35). The rate constants for reductive dissolution of 6-Mn0, by substituted anilines were correlated with Hammett constants u and u+. A better correlation was obtained with u+ values, for which Figure 14 shows a general trend between the rate constant and g+. That the reaction series followed u+ better than CT suggests extensive resonance interaction in the transition state (21). This is consistent with the subsequent discussion, which suggests the formation of a cation-radical intermediate. Nitroaniline does not follow the sequence and may interact with 6-Mn02by a different mechanism, as suggested later. Substituent effects and their correlations with physical and chemical properties provide a rational basis for estimating trends in the relative reactivity of para-substituted phenols and anilines with manganese dioxide. The order of the reactivity with substitution of the parent molecule is as follows: methoxy >> methyl > chloro > carboxy >> nitro. Those substituents that are considered classical electron-withdrawing groups make reductive dissolution of manganese dioxide by phenol or aniline more difficult by decreasing electron density, while the opposite effect on dissolution rate occurs for classical electron-donating groups. Similar orders of relative reactivity have been
azobenzene
Flgure 15. Postulated mechanism for oxidative coupling of aniline by reaction with manganese dioxide. The reaction proceeds from a cation radical through coupling products, which then undergo further oxidation [adapted from Sharma et al. (37)].
obtained for the oxidation of substituted anilines in dilute acetic acid medium by thallium triacetate and sodium iodate (16, 17). Oxidative Coupling of Anilines. The oxidation of aniline and substituted primary aromatic amines can result in the formation of a variety of products, depending on reaction conditions. The oxidation of aniline by MnOz in sulfuric acid gives p-benzoquinone (22), while oxidation with peroxy acids, e.g., peroxyacetic acid, or nitric acid may yield nitroso and nitro compounds (23, 24). Polymeric substances result from the reaction of aniline and peroxydisulfate in acid medium (25). Azo compounds result from the oxidation of primary aromatic amines by lead tetraacetate in acetic acid (26), by phenyl iodosoacetate in acetic acid (27), by N-chlorobenzamide in methanol/ hydrochloric acid medium (28),and by cuprous chloride and oxygen when pyridine is used as solvent (29). Substituted anilines can be oxidized to symmetrically substituted azobenzenes by refluxing with manganese dioxide in benzene (30). The results of this work show that the principal oxidation products of aniline and p-toluidine by manganese dioxide in aqueous buffered solution at pH 4 were azobenzene and 4,4'-dimethylazobenzene, respectively. A postulated reaction sequence may be inferred from cyclic voltammetry studies of primary aromatic amines in acidic aqueous solution as summarized in Figure 15 from the work of Sharma et al. (31), who examined the anodic oxidation of aniline, 2,4- and 2,6-dimethylaniline, and toluidines. According to this oxidation mechanism, aniline loses one electron to produce a cation radical, which through resonance structures may undergo head-to-tail, tail-to-tail, or head-to-head couplings. The immediate respective products are p-aminodiphenylamine,benzidine, and hydrazobenzene, all of which are more readily oxidized than the parent substance and undergo further oxidation at the same potential. The overall oxidation is thus a sequential, two-electron reaction. A suggested reaction pathway for p-toluidine is shown in Figure 16 in which head-to-head coupling is favored because of blocking at the para position. The results of this study confirm that the principal pathway for oxidation of aniline and p-toluidine by manganese dioxide in aqueous suspension at moderate acidity is a two-electron process, which results in the production of symmetrical azobenzenes. Head-to-head coupling with Environ. Sci. Technol., Voi. 24, No. 3, 1990 371
NH,
.*
N-N-H
2 e-
2eAzobenzene
CH3 p - Toluidine
H 3 Co y ; H
+
; l o c H3C-.. H
c-3
H,c - @- X e C H 3 + 2 H+ 4,4'- Dimethylhydrazobenzene
4,4'- Dlmethylozobenzene
Flgure 16. Postulated reaction mechanism for oxidative coupling of p -toluidine by reaction with manganese dioxide resulting in 4,4'dimethylazobenzene as the principal reaction product [adapted from Sharma et al. (37)].
formation of azo compounds is the predominant reaction, although a variety of minor products, such as p-aminodiphenylamine, the aniline head-to-tail intermediate coupling product, are evident also from reaction with manganese dioxide. These results are consistent with the work of Wheeler and Gonzalez (30),who demonstrated that manganese dioxide in benzene solution oxidizes substituted anilines to symmetrically substituted azobenzenes. They observed that nitroanilines and p-aminobenzoic acid were essentially unreactive over 24 h of refluxing in benzene with Mn02. The lack of reactivity of nitro- and carboxy-substituted aromatic amines was attributed to the nitro or carboxyl group being preferentially adsorbed at the active oxidation centers on the manganese dioxide, thus preventing oxidation of the amino group. Although the exact nature of the mechanisms involved in the oxidation of the anilines was not explored in the investigation reported here, the general features of the reaction with regard to oxidation products and relative rates of reaction agree with the findings of Wheeler and Gonzalez (30)and Sharma et al. (31). Additional studies with multiple-substituted anilines could help elucidate the reaction mechanism with MnOz in aqueous medium. For example, studies with 3,5-dichloro-4-nitroaniline or 3,5-diethyl-4-nitroaniline, in which adsorption of the nitro group would be hindered by the neighboring groups, would indicate the degree to which adsorption of the nitro group impedes the reaction.
Conclusion Transformation of Aniline and Azo Compounds. The present investigation demonstrated a redox pathway for oxidation of aniline and primary substituted anilines to symmetrical azo compounds by reaction with manganese dioxide in acidic aqueous systems. This oxidation reaction may be an important transformation process for aniline compounds in moderately acidic subsurface environments in the absence of molecular oxygen and substantial microbial activity. Analogous abiotic reductive transformations of azobenzene and selected azo derivatives have been observed by Weber and Wolfe (32) in sedimentlwater systems. The abiotic reduction of azobenzene was a four-electron, pH-dependent process for which GC-MS analyses showed production of stoichiometric amounts of aniline. The proposed scheme for the abiotic reduction of azobenzene to aniline is illustrated in Figure 17. The reaction is postulated to be a surface-mediated reaction proceeding through hydrazobenzene, which would not be stable in a highly reducing environment. The present investigation suggests that the pattern of oxidation of Environ. Sci. Technol., Vol. 24, No. 3, 1990
Aniline
Figure 17. Proposed scheme for abiotic reduction of azobenzene to aniline in anaerobic sediment/water systems [Weber and Wolfe (32)].
Cation radical
I
372
H ydrazobenzene
aniline by manganese dioxide is the reverse of the reduction of azobenzene in anaerobic sedimentlwater systems. The significance of the processes whereby manganese dioxide induced oxidative coupling of aniline and substituted anilines may serve as an elimination pathway for anilines depends on the accessibility and the form of manganese dioxide in natural systems, as well as the role and relative rates of competing reactions. Aniline may undergo microbial degradation through dioxygenase attack, resulting in oxidative deamination to catechol, which is further degraded by ortho cleavage and mineralized (7,33). Substituted anilines can be relatively resistant to microbial degradation, depending on the microbial community and the presence of other growth substrates (6, 10, 34). The abiotic redox reaction of primary amines and azo compounds in mineralogical and sedimentlwater environments may result in various oxidative-coupling and reductive-decoupling reactions. These processes may be significant with regard to the persistence and transformation of these classes of organic contaminants in environmental systems.
Acknowledgments
S. Shanmuganantha assisted with a portion of the experimental studies. Curt M. White and Louise J. Douglas, Coal Science Division, Pittsburgh Energy Technology Center, assisted with identification of reaction products. Registry No. 6-Mn02,1313-13-9;hydroquinone, 123-31-9; methyl-HQ,95-71-6;2,3-dimethyl-HQ,608-43-5;tert-butyl-HQ, 1948-33-0; trimethyl-HQ, 700-13-0; catechol, 120-80-9; 4methylcatechol, 452-86-8; p-anisidine, 104-94-9; p-toluidine, 106-49-0; aniline, 62-53-3;p-chloroaniline, 106-47-8; aminobenzoic acid, 1321-11-5;p-nitroaniline,100-01-6;azobenzene, 103-33-3; 4-aminodiphenylamine,101-54-2;1-methyl-1-imidazole,616-47-7; quinoline, 91-22-5; 5,5-dimethylhydantoin, 77-71-4; 4,4'-dimethylazobenzene, 501-60-0. Literature Cited (1) Stone, A. T.; Morgan, J. J. Enuiron. Sci. Technol. 1984,18, 450. (2) Stone, A. T.; Morgan, J. J. Enuiron. Sci. Technol. 1984,18, 617. (3) Stone, A. T. Enuiron. Sci. Technol. 1987, 21, 979. (4) Murray, J. W. J. Colloid Interface Sci. 1974, 46, 357. (5) Harrison, F. L.; Mallon, B. Selection of Representative Organic Compounds for in-depth Studies of Mobility in Soil- Water Systems; Lawrence Livermore National Laboratory: DE84-011625; UCID-20031,NTIS, 1982. (6) Lyons, C. D.; Katz, S. E.; Bartha, R. Bull. Enuiron. Contam. Toxicol. 1985, 35, 704. (7) Hwang, H.; Hodson, R. E.; Lee, R. F. Appl. Enuiron. Microbiol. 1985, 50, 1177. (8) Doull, J.; Klaassen, C. D.; Amdur, M. 0. Toxicology,2nd ed.; Macmillan Publishing Co.: New York, 1980. (9) Sax, N. I.; Lewis, R. J. Hazardous Chemicals Desk Reference; Van Nostrand Reinhold: New York, 1987. (10) Lyons, C. D.; Katz, S. E.; Bartha, R. Bull. Enuiron. Contam. Tonicol. 1985, 35, 697. (11) Lyons, C. D.; Katz, S. E.; Bartha, R. Appl. Enuiron. Microbiol. 1984, 48, 491. (12) Zepp, R. G.; Schlotzhauer, P. F. Enuiron. Sci. Technol. 1983, 17, 462. (13) Eight Peak Index of Mass Spectra, 1st ed.; Mass Spectrometry Data Centre: Reading, UK, 1970.
Environ. Sci. Technol. 1990, 24,373-378
Hites, R. A. CRC Handbook of Mass Spectra of Environmental Contaminants; CRC Press: Boca Raton, FL, 1985. Srivastava, S. P.; Gupta, R. C.; Shukla, A. K. Indian J . Chem. 1977,15A, 605. Radhakrishnamurti,P. S.; Pati, S. N. Indian J. Chem. 1980, 19A, 980. Rao, M. D. P.; Padmanbha, J. Indian J. Chem. 1980,19A, 984. Suatoni,J. C.; Snyder,R. E.; Clark, R. 0. Anal. Chem. 1961, 33, 1894. Bard, A. J.; Faulkner, L. R. Electrochemical Methods; Wiley: New York, 1980. Lyman, W. J.; Reehl, W. F.; Rosenblatt, D. H. Handbook of Chemical Property Estimation Methods; McGraw-Hill Book Co: New York, 1982. March, J. Advanced Organic Chemistry, 3rd ed.; Wiley: New York, 1985. LeNoble, W. J. Highlights of Organic Chemistry-An Advanced Text book; W. B. Saunders: Philadelphia, PA, 1965.
Plesnicar, B. In Oxidation i n Organic Chemistry; Trahanovsky, W. S., Ed.; Academic Press: New York, 1973; Vol. 5-C, pp 267-270. Ogata,Y. In Oxidation in Organic Chemistry; Trahanovsky, W. S., Ed.; Academic Press: New York, 1973; Vol. 5-C, pp 328-330.
Gupta, R. C.; Srivastava, S. P. Indian J. Chem. 1972,10, 706.
Pausacker, K. H.; Scroggie, J. G. J. Chem. SOC.1954,4003, part 11.
(27) Barlin, G. B.; Pausacker, K. H.; Riggs, N. V. J. Chem. SOC. 1954, 3122, part 111. (28) Rawat, B. S.; Agrawal, M. C. Indian J . Chem. 1978,17A, 299. (29) Nigh, W. G. In Oxidation in Organic Chemistry; Trahanovsky, W. S., Ed.; Academic Press: New York, 1973; Vol. 5-B, pp 51-53. (30) Wheeler, 0. H.; Gonzalez, D. Tetrahedron 1964,20, 189. (31) Sharma, L. R.; Manchanda, A. K.; Singh, G.; Verma, R. S. Electrochem. Acta 1982, 27, 223. (32) Weber, E. J.; Wolfe, N. L. Environ. Toxicol. Chem. 1987, 6 , 911. (33) Scow, K. M.; Simkins, S.; Alexander, M. Appl. Environ. Microbiol. 1986, 51, 1028. (34) You, I. S.; Bartha, R. J. Agric. Food Chem. 1982,30,274. (35) Hammett, L. P. Physical Organic Chemistry, 2nd ed; McGraw-Hill Book Co.: New York, 1970. (36) Bollag, J.-M.; Minard, R. D.; Liu, S.-Y. Enuiron. Sci. Technol. 1983, 17, 80-83. (37) Sillen, L. G.; Martell, A. E. Stability Constants for Metal-Ion Complexes; The Chemical Society: London, 1971; Supplement No. 1.
(38) Smith, R. M.; Martell, A. E. Critical Stability Constants;
Plenum Press: New York, 1975; Vol. 2. Received for review December 28, 1988. Revised manuscript received August 20,1989. Accepted November 27,1989. This work was supported by the U S . Department of Energy, Grant DE-FG22-86PC90524; Richard P. Noceti, Technical Project Officer,provided useful comments.
Anaerobic Microbial Remobilization of Toxic Metals Coprecipitated with Iron Oxide Aroklasamy J. Francls" and Cleveland J. Dodge Department of Applied Science, Brookhaven National Laboratory, Upton, New York 11973
An anaerobic N2-fixing Clostridium sp. solubilized Cd, Cr, Ni, Pb, and Zn coprecipitated with goethite (a-FeOOH) by (i) direct action due to enzymatic reduction of ferric iron and the release of metals associated with iron and (ii) indirect action due to metabolic products. The extent of dissolution depended upon the nature of the association of the metals with goethite. Substantial amounts of Cd and Zn, which were closely associated with iron, were released due to direct action. Nickel was solubilized by direct and indirect actions, while a small amount of Cr was solubilized only by direct action. The nature of association of P b in the coprecipitate was not affected by the presence of other cations and it was solubilized by indirect action. In the presence of bacteria, the concentration of soluble P b decreased due to biosorption. These results show that there could be significant remobilization of toxic metals coprecipitated with iron oxides in wastes, contaminated soils, and sediments due to microbial iron reduction.
Introduction Iron oxides scavenge transition and heavy metals in soils, sediments, and energy wastes (1-8). These oxides are a major sink for metals in the terrestrial and aquatic environments and play an important role in regulating their availability. Sorption and coprecipitation are the predominant processes by which most of the metals are retained by iron oxide. Sorption is a process by which metals are bound to the surface of an existing solid by adsorption and surface precipitation (9),whereas coprecipitation is 00 13-936X/90/0924-0373$02.50/0
the simultaneous precipitation of a chemical element with other elements and includes mixed-solid formation, adsorption, and inclusion ( 1 0 , I I ) . Although coprecipitation has not been as well studied as adsorption, it appears to remove trace metals from solution more efficiently (12-14). Toxic metals such as As, Cd, Co, Cr, Cu, Hg, Ni, Pb, Se, U, and Zn from fossil- and nuclear-fuel cycle waste streams, geothermal fluids, and electroplating wastes are currently removed by coprecipitation with ferric iron or are under consideration for such treatment (3,15-17). The removal of toxic metals from waste streams by coprecipitation with iron seems to be a very efficient method and economically feasible, but problems remain with the disposal of the solids generated in the process and with the ultimate fate of the coprecipitated metals in the environment. Significant dissolution of metals from the coprecipitate can be brought about by chemical and microbiological action. In general, solubility of iron oxide depends upon the degree of crystallinity. Amorphous iron oxides are orders of magnitude more soluble than goethite or hematite ( I ) , while amorphous synthetic goethite is 2-100 times more soluble than a well-crystallized goethite (18). Microorganisms play a significant role in the dissolution of amorphous and crystalline forms of iron oxides by direct action or by indirect action. Direct action involves enzymatic reductive dissolution of iron from higher oxidation state to lower oxidation state and indirect action is due to the production of metabolites (19-23). However, we have little information on microbially mediated disso-
0 1990 American Chemical Sc)ciety
Environ. Sci. Technol., Vol. 24, No. 3, 1990 373
